Catalysts & Activation Energy

The energy of a reaction with and without a catalyst present to lower activation energy

What is a catalyst?

Let’s talk about what catalysts are. A catalyst is a compound or element that increases the rate of a chemical reaction, e.g. the speed at which it occurs, without itself being part of the reaction. Generally speaking, a catalyst is not destroyed, consumed, or permanently changed in the reaction. A catalyst does this by lowering the activation energy, which we explain in the next section.

Catalysts work by providing an alternative pathway for the reaction to follow that has that lower activation energy, which makes it easier for the reactants to form the products. This allows the reaction to occur more quickly and efficiently. Catalysts can be either natural or synthetic, and they are commonly used in industry to make chemical reactions more efficient and cost-effective.

For example, enzymes are natural catalysts that are essential for many of the chemical reactions that take place in our bodies, and industrial catalysts are used in the production of a wide range of products, including plastics, fuels, and pharmaceuticals.

What is activation energy?

In simple terms, the definition of activation energy is the energy needed to start a reaction between two or more elements or compounds. A reaction with high activation energy may proceed slowly, or not at all. To speed up or start a reaction, you can either add the necessary activation energy or use a catalyst which lowers this requirement, effectively speeding up the reaction.

Activation energy units are in terms of energy units. The most common units used are KJ/mol or J/mol.

How does a catalyst work?

A catalyst increases the rate of reaction by decreasing the activation energy.  Decreased activation energy means less energy required to start the reaction.

The graph below shows the energy of a reaction both with and without a catalyst present. The x-axis is the reaction coordinate or progression of the reaction from reactant (left side) to product (right side). The y-axis is the energy.

Activation energy of reaction with and without a catalyst. Catalysts lower the activation energy to increase the rate of reaction.
The energy of a reaction with and without a catalyst. (Source: Wikipedia Commons)

With the catalyst present, the activation energy (Ea) is smaller. Visually, the hill the reaction has to climb before going downhill to the products is smaller. Just as riding a bike over a small uphill is easier than a larger uphill, a reaction proceeds faster when the activation energy hill is smaller.

A catalyst lowers the activation energy by changing the transition state of the reaction. The reaction then goes through a different pathway/mechanism than the uncatalyzed reaction. The catalyst does not change the net energy difference between reactant and product.  The reaction’s net equation will be the same in a catalyzed and uncatalyzed reaction even though the transition state changes.

Overall reaction: A + B + catalyst –> AB + catalyst

Net Reaction: A + B –> AB

Main Categories of Catalysts

Heterogeneous Catalysts

A heterogeneous catalyst is in a different phase than the reactants. Usually, that means the catalyst is in the solid phase and reactants are in the liquid or gas phase. Another name for a heterogeneous catalyst is a surface catalyst.

Heterogeneous catalysts work by attaching the catalyst to a solid support structure and the reactants flow over and past the catalyst, reacting along the way. A benefit of this type of catalyst is that the catalyst is easily separated from the product when the reaction is complete. The catalyst can then easily be reused. In manufacturing, this is an important cost-cutting measure. A drawback of the heterogeneous catalyst is that the amount of interaction between reactant and catalyst can be limited by surface area and diffusion of the product away from the surface.

A common heterogeneous catalyst is a catalytic converter for gasoline in cars. Another important heterogeneous catalyst is the Haber-Bosch process which forms NH3.  

Homogeneous Catalysts

In a homogeneous catalyst, both the reactants and the catalyst are in the same phase. Normally they are both in either the liquid or gas phase.

The main benefit of a homogeneous catalyst is the increased interaction between reactant and catalyst. Both can move freely and are therefore more likely to interact and lead to a reaction.

Common homogeneous catalysts are transition metals and acids. One homogeneous catalyzed reaction is the conversion of oxygen to ozone in the atmosphere. Nitric oxide (NO) catalyzes the reaction. All of the participants in the reaction reside in the gas phase. Therefore we know it is a homogeneous catalytic reaction.


Enzymes are large proteins that are biological catalysts. They are powerful forces in the body. Often they catalyze only one very specific reaction (compared to inorganic catalysts that often catalyze a much more broad set of reactions). The specificity is due to the active site in the catalyst—a pocket of specific chemical composition formed by amino acids where only one very specific reactant model will fit. This is also referred to as the lock-and-key model.

Enzymes play a lot of important roles in the body. They catalyze the breakdown of starch to create glucose. They also convert carbon dioxide (CO2) to other molecules the body needs such as HCO3. Enzymes assist and sped up almost all processes in the body.

Calculating Activation Energy

Activation energy calculations use the Arrhenius equation. We will cover the basics here, but for more examples and in-depth analysis be sure to check out the Arrhenius equation article.

The Arrhenius equation is:

k = A \cdot e^{-\frac{E_a}{RT}}

Where k is rate constant, E_a is the activation energy, A is the frequency factor, R is the gas constant, and T is temperature.

We can find the activation energy if we know the rate constant (k) at various temperatures (T). To determine the activation energy (E_a) we plot ln(k) vs 1/T. By doing so we get a line with a slope of –E_a/R and a y-intercept of A. R is a constant, so then we can solve for activation energy.

For example problems on the activation energy formula, see the Arrhenius equation page!

List of common catalysts:

  • Vanadium pentoxide (making sulfuric acid)
  • Palladium metal
  • Manganese dioxide
  • Platinum metal (in catalytic converters)
  • Iron metal (in the Haber process)
  • Aluminum chloride (many organic reactions)
  • Copper (II) oxide
  • Enzymes (naturally occurring catalysts in biochemical reactions)

Catalysts & Activation Energy Vocabulary Definitions

  • Activation Energy – energy needed to start a reaction between two or more elements or compounds
  • Catalyst– A molecule that increases the rate of reaction and not consumed in the reaction
  • Turnover Number – the number of reactions one enzyme can catalyze per second
  • Enzyme – a biological catalyst made of amino acids.
  • Lock-and-key model- The model that an enzyme and the reactant molecule have a similar shape at the active site to increase specificity and efficiency of the reaction.
  • Homogeneous Catalyst – A catalyst present in the same phase as the reactants.
  • Heterogeneous Catalyst- A catalyst in a different phase than the reactants.

We hope you now know what is a catalyst, how they effect a reaction, and how the activation energy comes into play. These concepts will become more important if you delve more deeply into kinetics of chemical reactions. Other factors can also influence reaction rate, to learn more about those, continue reading here!